For which one of the following reactions will there be a positive $∆\mathrm{S}$?

• ${\mathrm{H}}_2\mathrm{O}\left(\mathrm{g}\right) \to {\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$

• ${\mathrm{H}}_2 + {\mathrm{I}}_2 \to 2\mathrm{HI}$

• ${\mathrm{CaCO}}_3\left(\mathrm{s}\right) \to \mathrm{CaO}\left(\mathrm{s}\right) + {\mathrm{CO}}_2\left(\mathrm{g}\right)$

• ${\mathrm{N}}_2\left(\mathrm{g}\right) + 3{\mathrm{H}}_2\left(\mathrm{g}\right)\to 2{\mathrm{NH}}_3\left(\mathrm{g}\right)$

The standard Gibb's free energy change, $∆$G° is related to equilibrium constant, K_p as

• ${\mathrm{K}}_{\mathrm{p}}= -\mathrm{RT} \ln ∆\mathrm{G}°$

• ${\mathrm{K}}_{\mathrm{p}}= {\left[\frac{\mathrm{e}}{\mathrm{RT}}\right]}^{∆\mathrm{G}°}$

• ${\mathrm{K}}_{\mathrm{p}}=-\frac{∆\mathrm{G}°}{\mathrm{RT}}$

• ${\mathrm{K}}_{\mathrm{p}}={\mathrm{e}}^{-\frac{∆\mathrm{G}°}{\mathrm{RT}}}$

The yield of the product in the reaction

${\mathrm{A}}_2\left(\mathrm{g}\right) + 2\mathrm{B}\left(\mathrm{g}\right) \rightleftharpoons \mathrm{C}\left(\mathrm{g}\right) + \mathrm{Q} \mathrm{kJ} \mathrm{would} \mathrm{be} \mathrm{higher} \mathrm{at}$

• high temperature and high pressure

• high temperature and low pressure

• low temperature and high pressure

• low temperature and low pressure

Given that $∆{\mathrm{H}}_{\mathrm{f}}$(H) = 218 kJ / mol, express the H-H bond energy in kcal/mol.

• 52.15

• 911

• 104

• 52153

365.

$$\begin{gathered} \mathrm{Calculate} ∆\mathrm{H}° \mathrm{for} \mathrm{the} \mathrm{reaction}, \\ {\mathrm{Na}}_2\mathrm{O}\left(\mathrm{s}\right) + {\mathrm{SO}}_3\left(\mathrm{g}\right) \to {\mathrm{Na}}_2{\mathrm{SO}}_4\left(\mathrm{g}\right) \\ \mathrm{given} \mathrm{the} \mathrm{following}: \\ \left(\mathrm{A}\right) \mathrm{Na}\left(\mathrm{s}\right) +{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) \to \mathrm{NaOH}\left(\mathrm{s}\right) + \frac{1}{2}{\mathrm{H}}_2\left(\mathrm{g}\right) ;∆\mathrm{H}°= -146 \mathrm{kJ} \\ \left(\mathrm{B}\right) {\mathrm{Na}}_2{\mathrm{SO}}_4\left(\mathrm{s}\right) +{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right) \to 2 \mathrm{NaOH}\left(\mathrm{s}\right) + {\mathrm{SO}}_3\left(\mathrm{g}\right); ∆\mathrm{H}°=+ 418 \mathrm{kJ} \\ \left(\mathrm{C}\right) 2 {\mathrm{Na}}_2\mathrm{O}\left(\mathrm{s}\right) +2 {\mathrm{H}}_2\left(\mathrm{g}\right) \to 4\mathrm{Na}\left(\mathrm{s}\right) + 2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right); ∆\mathrm{H}°= +259 \mathrm{kJ} \end{gathered}$$

• +823 kJ

• -581 kJ

• -435 kJ

• +531 kJ

If an endothermic reaction occurs spontaneously at constant temperature (T) and pressure (p), then which of the following is true?

• $∆\mathrm{G} > 0$

• $∆\mathrm{H} < 0$

• $∆\mathrm{S} > 0$

• $∆S < 0$

A spontaneous process is one in which the system suffer

• no energy change

• a lowering of free energy

• a lowering of entropy

• an increase in internal energy

When a solid melts reversibly :

• H decreases

• G increases

• E decreases

• S increases

Enthalpy is equal to :

• ${\mathrm{T}}^2{\left[\frac{\partial (\mathrm{G}/\mathrm{T})}{\partial \mathrm{T}}\right]}_P$

• $-{\mathrm{T}}^2{\left[\frac{\partial (\mathrm{G}/\mathrm{T})}{\partial \mathrm{T}}\right]}_{\mathrm{P}}$

• ${\mathrm{T}}^2{\left[\frac{\partial (\mathrm{G}/\mathrm{T})}{\partial \mathrm{T}}\right]}_V$

• $-{\mathrm{T}}^2{\left[\frac{\partial (\mathrm{G}/\mathrm{T})}{\partial \mathrm{T}}\right]}_{\mathrm{V}}$

Condition for spontaneity in an isothermal process is

• $∆\mathrm{A} + \mathrm{W} < 0$

• $∆\mathrm{G} + \mathrm{U} <0$

• $∆\mathrm{A} + \mathrm{U}>0$

• $∆\mathrm{G} -\mathrm{U}<0$